Synthesis, Structural Characterization and Biological Activity of Novel Cyclohexane-1,3-dione Ligands and Their Metal Complexes

Some new Zn(II) and Cu(II) complexes [Cu(L1)(OAc)2]∙H2O, [Cu(L1)(NO3)H2O]∙NO3∙3.5H2O, [Zn(L1)(NO3)2]∙4.5H2O, [Zn(L1)(OAc)2(H2O)2]∙3H2O, [Cu2(L2)(OAc)4]∙2H2O∙2DMF, [Cu(L2)2]∙2NO3∙1.5DMF∙H2O, [Zn(L2)2(NO3)2]∙DMF and [Zn2(L2)(OAc)4(H2O)4]∙5H2O; L1 = 2-[2-(2-methoxyphenyl)hydrazono]cyclohexane-1,3-dione and L2 = 2-[2-(3-nitrophenyl)hydrazono]cyclohexane-1,3-dione were synthesized and characterized by IR, 1H-NMR,13C-NMR and ultraviolet (UV-Vis.) spectroscopy, elemental analysis, magnetic susceptibility, mass spectrometry and thermogravimetry-differential thermal analysis (TGA-DTA). The synthesized ligands and their complexes were tested for antibacterial activity against Escherichia coli ATCC 25922, Enterococcus faecalis ATCC 29212, Staphylococcus aureus ATCC 25923, and Salmonella typhimurium CCM 583. Some of complexes showed medium-level antibacterial activity against the test bacteria compared with ampicillin.


Introduction
Cyclohexane-1,3-dione and its derivatives are important building blocks. The cyclohexane-1,3dione skeleton is a characteristic molecular fragment common for a class of natural and synthetic herbicides and drugs inhibiting 4-hydroxyphenylpyruvate deoxygenase [1,2]. The inhibitory properties of this class of compounds are a result of their ability to chelate the ferrous ion in the active site of the enzyme.
Arylhydrazones of mono-, di-and triketones as well as cyclic 1,3-diones viz., cyclohexane-1,3dione have been reported in the literature. Such compounds have been extensively used as precursors of potential anti-diabetic drugs [3]. Hydrazones are compounds whose molecules contain the C=N-NH-triatomic linkage. The biological, chemical and industrial versatility of hydrazones and their complexes continue to attract considerable attention [4,5]. The role of hydrazones in treating tuberculosis is well known. Substituted hydrazones have established spasmolytic activity, activity against leukemia, and sarcomas [6]. Also they possess anti-inflammatory, analgesic, antipyretic, antibacterial and antitumor activities [7].

Infrared Spectra
IR absorptions of both L 1 and L 2 and their metal complexes along with assignments are summarized in the experimental section. All complexes are air stable and insoluble in most organic solvents and water, but freely soluble in coordinating solvents such as pyridine, DMF or DMSO. All have higher melting or decomposition points than the parent ligands.
In the IR spectrum of L 1 the characteristic peaks are seen at 2955-3500 cm −1 ν(NH), 1280 cm −1 ν(OCH3), 1696, 1671 cm −1 ν(C=O), and 1505 cm −1 ν(C=C) [8,[9][10][11]. The L 1 shows a broad NH band at 3343 cm −1 due to the hydrazone NH group. The position and broadness of the band are indicative of intramolecular hydrogen bonding between the hydrazone proton and the carbonyl oxygen. The hydrazone can coordinate to transition metals either in the enolic form (N=N-C=C-OH) or in ketonic form (HN-N=C-C=O). In L 1 and L 2 the characteristic hydrazone group >CONH-peaks indicate that the ligand exists in the keto form in the solid state. This NH band shifted upon complexation of the ligand with metals due to coordination of the hydrazone proton in [Cu2(L 2 )(OAc)4]·2H2O·2DMF and all of the Zn(II) complexes [12,13].
In the IR spectrum of L 2 , the characteristic peaks are at 1686 and 1353 cm −1 , which are assigned to ν(C=O) and ν(NO2), and at 1614 cm −1 which is assigned to the ν(C=N) group [15,16]. In addition, the spectrum of the ligand L 2 shows broad bands in the 3500-3100 cm −1 region, which may be assigned to a ν(NH) band. In the IR spectra of L 2 , the bands assigned to ν(C=O) and ν(C=N) are shifted by ±16-35 cm −1 in the spectra of complexes, indicating coordination through the exocyclic carbonyl oxygen and C=N nitrogen of L 2 . In the spectra of the metal complexes of the L 2 , the band at 1353 cm −1 for the NO2 group did not shift [17], showing that in these complexes the NO2 group did not participate in complex formation. In the IR spectra of the complexes, ν(OCH3) remains unmodified, indicating that the methoxy group is also not involved in the coordination.
The infrared spectra of all complexes, except that of [Zn(L 2 )2(NO3)2]·DMF complex, exhibited intense broad bands at 3560-3311 cm −1 that are attributed to ν(OH) of the lattice or coordinated water molecules, while the ν(H2O) bands observed at approximately 876-764 cm −1 are assigned to coordinated water molecules [18][19][20] The ν(N-N) bands of the L 1 and L 2 ligands at 1025, 1010 cm −1 are found to be shifted to higher energies (1041-1020 cm −1 ) in the spectra of the complexes, indicating coordination via the nitrogen atoms. This is confirmed by bands in the range of 491-455 cm −1 , which have been assigned to the ν(Zn-N) ν(Cu-N) bands [21].

1 H-NMR and 13 C-NMR Spectra
The 1 H-NMR spectra of L 1 , L 2 and their Zn(II) complexes were recorded in dimethylsulfoxide (DMSO-d6) solution using tetramethylsilane (TMS) as internal standard. The 1 H-NMR spectrum of L 1 shows a singlet at 4.00 ppm that may be assigned to the protons of the methoxy group, a singlet at 15.00 ppm that may be assigned to the proton of the NH group, a multiplet at 7.20-7.80 ppm that may be assigned to the aromatic protons and another at 2.00-2.70 ppm that may be assigned to the hexane ring protons. Similar peaks are observed for the ligand L 2 . The 1 H-NMR spectrum of the diamagnetic Zn(II) complexes shows almost the same values as the ligands. The signals due to the -NH proton are shifted downfield in the spectra of the Zn(II) complexes, indicating the coordination of the ligand through the nitrogen of the -NH groups to the metal ions. The hydrogen bonding decreases the electron density around the proton and thus moves the proton absorption to a lower field [27]. In addition, H2O and acetate proton signals are seen in NMR spectra of complexes.

Electronic Spectra and Magnetic Properties
The electronic spectra of L 1 , L 2 and their metal complexes were recorded in 3 × 10 -3 molar DMF solutions in the range from 200 to 800 nm. The electronic spectrum of the ligands shows broad bands at 218, 242 and 316-393 nm; the first two absorption bands may be assigned to n→π* and π→π* transitions of >C=O and >C=N moieties [28], the third absorption band may be due to high-intensity charge transfer transitions.
The [Cu(L 1 )(OAc)2]·5H2O complex shows square-planar geometry. It shows intense absorption bands in the 254-269 and 414 nm range which can be assigned to the π→π* transition of the C=N or C=O groups and charge-transfer bands, respectively. The electronic spectrum of the complex shows an absorption band in the range 580 nm, attributed to 2 T2g→ 2 Eg transition suggesting a square-planar geometry [29].
Square planar geometry is suggested for [Cu2(L 2 )(OAc)4]·2H2O·2DMF from the presence of bands at 600 and 454 nm for the complex. These bands are assigned to the 2 B1g→ 2 A1g and 2 B1g→ 2 Eg transitions, while the second band is due to charge transfer. The magnetic moment values for the Cu(II) complex are normal and lower (1.66 B.M.) than expected 1.7-2.2 B.M. for one unpaired electron [32,33].
The electronic spectrum of Cu(II) complex shows bands in the regions of 254, 376 and 470 nm which may be assigned to the 2 B1g→ 2 A1g transitions in a square-planar geometry. The magnetic moment value of [Cu(L 2 )2]·2NO3·1.5DMF·H2O complex is 2.22 B.M. [34].
The electronic spectra of the Zn(II) complexes show only a high-intensity band at 357-425 cm −1 assigned to ligand-metal charge transfer.

Thermal Studies
The thermal stability of the ligands and complexes is investigated using thermogravimetric analysis. The TGA curves are obtained at a heating rate of 10 °C·min −1 under a nitrogen atmosphere in a temperature range of 25.0-800.0 °C. The results obtained are in good agreement with the theoretical formula suggested from the elemental analyses. The thermal data are summarized in Table 1.
The [Cu2(L 1 )2(AcO)2]·5H2O complex is stable up to 95 °C and its decomposition starts at this temperature. A 16.92(17.39%) weight loss is observed at 195 °C corresponding to five moles of water of crystallization. The DTA curve of the complex shows one endothermic peak at 350.50-360.00 °C. For the Cu(II) complex, the DTA curve shows an endothermic process around 350.50 °C that corresponds to the rupture of the coordinated bond and simultaneous melting of the complex [35].
The [Cu(L 1 )(NO3)H2O]·NO3·3.5H2O complex was stable up to 85 °C and its decomposition started at this temperature. The first decomposition stage at 160 °C is attributed to the removal of 3.5 moles of water of crystallization and 1 mole of nitrate ion with a 24.62(24.29%) loss. The second weight loss stage of 234.48-494.55 °C corresponds to the removal of the H2O, NO3and C6H6O2N as a fragment of the ligand decomposition. The DTA curve of the Cu(II) complex shows two endothermic peaks at 150.10 °C, 380.50 °C and one exothermic peak at 485.25 °C.
The [Zn(L 1 )(NO3)2]·4.5H2O complex decomposed in three stages. In the first stage, loss of 4.5 moles of water of crystallization and in the second stage, loss of two moles of NO3and C13H14O2N2 groups resulted. In the third stage, the complex decomposed to ZnO. The DTA curve of the complex shows two endothermic peaks at 450.00-470.20 °C. For the Zn(II) complex, these peaks corresponds to simultaneous melting of the complex [35].
The [Zn(L 1 )(OAc)2(H2O)2]·3H2O complex decomposed in two stages. The first is due to removal of the crystallization water molecule with a weight loss of 18.76(18.83%). The second weight loss is associated with the removal of two moles of water as well as 2C2H3O2 − from the ligand [36]. The DTA curve of the Zn(II) complex shows two endothermic peaks at 280.10 °C, 330.50 °C and one exothermic peak at 420.25 °C. [Cu2(L 2 )(OAc)4]·2H2O·2DMF shows two decomposition steps. The first decomposition step in the temperature range of 180.12-240.32 °C may be attributed to the loss of crystallization water molecules and two DMF molecules [36,37]. The second step within the temperature range of 240.32-285.34 °C 29.23(29.27%) corresponds to the removal of four moles of acetate molecules as deduced from weight loss calculations. The DTA curve of the Cu(II) complex shows two endothermic peaks at 160.50 °C and 190.10 °C. The endothermic peaks are likely due to the loss of water of hydration and DMF molecules.
The complex [Cu(L 2 )2]2NO3·1.5DMF·H2O is stable up to 100 °C and its decomposition started at this temperature. In the decomposition process of the Cu(II) complexes, the mass losses corresponded to one mole H2O, 1.5 moles DMF and two moles of NO3 − in the first stage of the decomposition, respectively. The Cu(II) complex is stable up to 160.00-240.00 °C and after this temperature other decomposition ensues. The DTA curve of the Cu(II) complex shows endothermic peaks at 100.20 °C and 215.50 °C corresponding to the loss of H2O, DMF and NO3and one exothermic peak at 500.10 °C corresponding to the second weight loss.
The [Zn2(L 2 )(OAc)4(H2O)4]·5H2O complex followed a four-staged decomposition, in which five moles of water of crystallization, four moles of coordinated water, four moles of acetate ions, and an 2-iminocyclohexane-1,3-dione group and the rest of the ligand decomposed successively. The DTA curve of the Zn(II) complex shows two endothermic peaks at 130.50 °C and 180.10 °C. The endothermic peaks are likely due to the loss of molecules of water of hydration.
The decomposition curve of [Zn(L 2 )2(NO3)2]·DMF begins with a step at 160.53-241.20 °C, displaying a 9.23(9.30%) weight loss corresponding to the removal of DMF [38]. The second step within the temperature range of 241.20-277.55 °C corresponds to the removal of two moles of NO3 − and C12H14N2O2 groups 44.61(44.10%) as deduced from weight loss calculations. The DTA curve of the Zn(II) complex shows one endothermic peak at 170.50 °C and two exothermic peaks at 300.60 °C and 350.40 °C. The first one is related to removal of the DMF molecule, whereas the other is the second step of decomposition [18].
The final products formed during thermal analysis of the complexes could not be determined because the decompositions of the complexes were not complete at 800 °C, except for the [Zn(L 1 )(NO3)2]·4.5H2O and [Cu(L 2 )2]·2NO3·1.5DMF·H2O complex.

Mass Spectra
The mass spectra peaks of the [Cu(L 1 )(OAc Single crystals of the complexes could not be isolated from any solutions, thus no definite structure could be described. However, the analytical, spectroscopic and magnetic data enable us to propose the possible structures shown below in Figures 2-9.

Biological Evaluation
The synthesized compounds were screened for their antibacterial activity using Escherichia coli ATCC 25922, Enterococcus faecalis ATCC 29212, Staphylococcus aureus ATCC 25923, and Salmonella typhimurium CCM 583. The results are shown in Table 2
Elemental analyses were carried out on a CHNS-O model 932 elemental analyzer (Leco, St. Joseph, MI, USA). 1 H-NMR and 13 C-NMR spectra were recorded using a model DPX-400 MHz FT spectrometer (Bruker GmbH, Billerica, MA, USA). IR spectra were recorded on a Precisely Spectrum One spectrometer (Perkin Elmer, Akron, OH, USA) using KBr discs in the wavenumber range of 4000-400 cm −1 .
Electronic spectra studies were conducted on a model UV-1700 spectrophotometer (Shimadzu, Kyoto, Japan) in the wavelength 1100-200 nm. Magnetic susceptibility measurements were performed using the standard Gouy tube technique. Hg[Co(SCN)4] was used for calibration calibrate. LC/MS-API-ES mass spectra were recorded using an Agilent model 1100 MSD mass spectrophotometer (Minneapolis, MN, USA). Thermogravimetric analysis (TGA) and differential thermal analysis (DTA) were carried out in nitrogen atmosphere with a heating rate of 10 °C·min −1 using a Shimadzu DTG-60 AH thermal analyzer.

Antibacterial Activity Studies
The disk diffusion method was used for determining the antibacterial activity of ligands and complexes. Antibacterial activity against Enterococcus faecalis ATCC 29212, Staphylococcus aureus ATCC 25923, Escherichia coli ATCC 25922, and Salmonella typhimurium CCM 583 was investigated. Mueller-Hinton agar (Oxoid Ltd., Basingstoke, Hampshire, UK) was used for all bacterial strains, except for Enterococcus faecalis ATCC 29212 for which Mueller Hinton agar with 5% defibrinated sheep blood was used. The media were melted at 100 °C, autoclaved at 121 °C for 15 min, cooled 45 °C to 50 °C and were poured into plates of 9 cm diameter in quantities of 20 mL, and left on a flat surface to solidify and the surface of media was dried at 37 °C. Then, preparation of the inoculums was used for colony growth method in Mueller-Hinton broth to a turbidity equivalent to a 0.5 McFarland (10 8 cfu·mL -1 ) standard. The organisms were streaked on Petri dishes using sterile cotton swab. The surface of the media was allowed to dry 3-5 min at room temperature. The 10 mg·mL -1 (in DMSO, Merck), of compound impregnated blank discs (Oxoid Ltd.) were applied to the surface of inoculated plates. The Mueller-Hinton agar plates were incubated at 35 ± 2 °C for 18-24 h. The plates were examined and the diameter of the inhibition zone was measured by surrounding discs. The antibiotic ampicillin (10 μg, Oxoid) was used as the standard [40,41].

Synthesis of the Ligands L 1 and L 2
A hydrochloric acid solution (2.5 mL) of 3-methoxyaniline (1.23 g, 10 mmol) and an aqueous solution (10 mL) of sodium nitrite (0.69 g, 10 mmol) were mixed and stirred at 273 K for 1 h. To this solution, an ethanol solution (10 mL) of the coupling component cyclohexane-1,3-dione (1.12 g, 10 mmol) was added and the stirring was continued at 273 K for 4 h. The resulting product was filtered and washed with water, dried and crystallized from ethanol (yield 77%). A similar synthesis was performed for L 2 using 3-nitroaniline (1.38 g, 10 mmol) instead of 3-methoxyaniline.

Conclusions
We have prepared two new ligands and their metal complexes. Structures of the ligands and complexes were confirmed by spectral and analytical techniques. Antibacterial activities of the complexes were tested against Escherichia coli ATCC 25922, Enterococcus faecalis ATCC 29212, Staphylococcus aureus ATCC 25923, and Salmonella typhimurium CCM 583, respectively. Some complexes showed average level antibacterial activity against the test bacteria compared to ampicillin, but some of the complexes did not exhibit antibacterial activity against the test microorganisms. For these reasons, these complexes may not be useful for preparing new active agents.